Identify each statement as true or false: (c) Cl^- is bigger than I^-.

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To determine the relative sizes of Cl^- and I^-, we need to consider the concept of ionic radii. Ionic radius is the measure of an atom's ion in a crystal lattice.

Recall that as you move down a group in the periodic table, the ionic radius increases. This is because additional electron shells are added, increasing the distance between the nucleus and the outermost electrons.

Chlorine (Cl) is in the third period, while iodine (I) is in the fifth period of the periodic table. Therefore, I^- has more electron shells than Cl^-.

Since I^- has more electron shells, it is larger in size compared to Cl^-. The increased number of electron shells outweighs the effect of increased nuclear charge.

Based on this understanding, the statement 'Cl^- is bigger than I^-' is false, as I^- is actually larger than Cl^-.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Ionic Radius

The ionic radius refers to the size of an ion in a crystal lattice. Cations (positively charged ions) are typically smaller than their neutral atoms due to the loss of electrons, while anions (negatively charged ions) are larger due to the gain of electrons. Understanding ionic radius is crucial for comparing the sizes of different ions.

Periodic Trends

Periodic trends are patterns observed in the properties of elements across the periodic table. For instance, ionic size generally increases down a group due to the addition of electron shells, while it decreases across a period due to increased nuclear charge. Recognizing these trends helps in predicting the relative sizes of ions.

Chloride vs. Iodide Ions

Chloride (Cl<sup>-</sup>) and iodide (I<sup>-</sup>) are both anions, but they differ in size. Iodide is larger than chloride because iodine is located further down in the periodic table, resulting in more electron shells. This difference in size is essential for understanding the statement in the question.

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